BOHR MODEL (approx 1913)¾The Bohr model of the atom was based on the line spectra of the

Hydrogen atom. The model also incorporated the concept developed by Einstein regarding the particle behaviour of light during emission or absorption (photon or quanta of energy).

Main Points:1. ENERGY LEVELS¾ an electron can only have specific energy values in an atomthe path followed by the electrons is a circular orbit (spherical in 3-D) these energies are called energy levels given the name Principal Quantum Number, n

¾ the orbit closest to the nucleus is given n=1, with the lowest energy as one moves outward the energies get larger and “n” increases (n=1,2,3,4,...)

¾ an electron can only circle in one of the allowed orbits WITHOUT a loss of energy

2.TRANSITION BETWEEN ENERGY LEVELS¾ an electron in an atom can only change energy by going from one energy

level to another level, i.e. NOT in-between

¾ light is emitted when an electron falls from a higher energy level to a lower energy level

¾ an electron generally remains in its ground state – the lowest energy level possible (n=1 for hydrogen)

¾when an electron is given energy it can be bumped to a higher energy level called the excited state

¾ as the electron drops from the excited state down to lower energy levels it emits light

¾Bohr was able to show that his model was able to match the Balmer series (Visible light, drop to n=2) and predict the Paschen series ( IR light , drop to n=3) and the Lyman series (UV light, drop to n=1)

¾Although the Bohr model could explain the behaviour of the hydrogen atom, it could not fully explain the behaviour of other atoms

The Bohr Model of the Atom

One major question with respect to the atom’s structure remained:

An orbiting electron emits radiation that should cause it to lose energy and collapse into the nucleus.

Bohr’s observations:

How Bohr explained the emission spectra and absorption spectrum of hydrogen gas:

Quantum Mechanics: Describing the Electron

The Quantum NumbersThese numbers are used to describe the, probable location of an electron.

i) The Principal Quantum Number (n)The value for the energy level or orbital of an e-1

(n= 1, 2, 3 to ∞).

ii) The Secondary Quantum Number (l)Divides the energy levels into sub-levels or orbital types (smaller spaces or shapes) called s (0), p (1), d

(2) and f (3). To determine the shape: l = 0 to n-1.

iii) The Magnetic Quantum Number (ml)Describes the direction of orientation (pointing) for p, d and f orbital types (ie. px, py, pz) To determine the direction: ml = -l to +l.

iv) The Spin Quantum Number (ms)An e-1 can spin in only 2 directions (ms = +½ or ms= - ½)

Spin Quantum Number, ms¾needed to explain additional spectral line-splitting & different kinds of magnetism

¾ferromagnetism-associated with substances containing Fe, Co & Ni

¾paramagnetism-weak attraction to strong magnets (individual atoms vs. collection of atoms)

¾To explain magnetism we can draw the electron configuration of a ferromagnetic element, e.g. Fe [Ar] 4s2

3d6 = 1 pair + 4 unpaired. It is the unpaired electrons cause the magnetism

¾paramagnetism couldn’t be explained until Wolfgang Pauli suggested that electrons spin on their axis (1925)

¾could spin only 2 ways (clockwise vs. counterclockwise) and he used only 2 numbers to describe this:

¾ ms = +1/2 (clockwise) or –1/2 (counterclockwise)

¾opposite pairs of electron spins represent a stablearrangement

¾when electrons are paired – spin in opposite directions –the magnetic field is neutralized, while an individual electron spin can be affected by a magnet

Quantum Numbers Questions

1) An electron has the following set of quantum numbers: n = 3, l = 1, ml = 1, ms = -½ .

In which orbital is this electron found?a. 3s b. 3p c. 3d d. 3f e. 4p

2) Which set of quantum numbers is not possible?a. n = 3, l = 0, ml = 0, ms = -½b. n = 5, l = 3, ml = 2, ms = ½c. n = 4, l = 3, ml = -1, ms = -½d. n = 5, l = 3, ml = -3, ms = ½e. n = 4, l = 4, ml = 2, ms = -½

3) Which set of quantum numbers is not possible?a. n = 5, l = 3, ml = 0, ms = -½b. n = 1, l = 0, ml = 0, ms = -½c. n = 3, l = 2, ml = 1, ms = ½d. n = 4, l = 3, ml = -3, ms = 0e. n = 5, l = 2, ml = 0, ms = ½

4) What are the allowed values of ml for an electron with each orbital-shape quantum number.a) l = 3 b) l = 1

5a) What are the possible values of ml if n = 4 and l = 2?

b) What kind of orbital is described by these quantum numbers?

c) How many orbitals can be described by these quantum numbers?

PROBABILITY MODEL

(WAVE MECHANICAL MODEL or QUANTUM MECHANICAL MODEL)DeBroglie (1924)

¾Proposed that moving electrons could be considered to behave like waves

¾experimentation showed that electrons can exhibit diffractionpatterns like light – a key part of the wave model of light derived expression relating wave properties to particle properties

ג = h/mv where: ג = wavelength of a particle wave

h = Planck’s Constantm = mass of particlev = speed of particlefor a baseball of m= 0.145 g with v=27 m/s

10= ג -34 mfor an electron of m = with v=3.00x108 m/s

100= ג x 10-12 m (100 pm)visible light has a wavelength range of 400 to 750 nm

Schrodinger’s Theory (1926)

¾extended wave properties of electrons to their behaviour in atoms and molecules

¾consider Bohr’s Model – electron orbits nucleus like the earth around the sun in a continuous path

¾in quantum mechanics – electron does NOT have a precise orbit in an atom

¾using principle of standing waves (wave theory) and the quantum numbers, Schrodinger was able to devise an expression, called the wave equation, that could identify the space an electron will occupy

Energy Level Diagrams and Electron ConfigurationsBoth methods are used to show the distribution of electrons among the various orbitals in an atom according to quantum mechanics. Rules to follow when drawing diagrams and writing electron configurations:

Aufbau (building up) PrincipleElectrons are added to the lowest energy orbital available.

Pauli Exclusion PrincipleAn orbital can be empty, have one electron or have two electrons (at most).

Hund’s RuleElectrons in the same sub-level will not pair up (occupy the same orbital) until all orbitals in the sub-level are half-filled (have 1 electron).

Energy Level Diagrams

Graphs the distribution of electrons according to energy and indicates direction of spin. e.g.i) Neon:

ii) Calcium +2

:

Electron Configurations

Electron configurations are a shorthand notation of energy level diagrams and provide information about the first two quantum numbers, n and l

Write electron configurations for the following:

i) Zinc

ii) Stable Calcium Cation

iii) Neon

iv) Mn

Memory Trick7s 7p

6s 6p 6d

5s 5p 5d 5f

4s 4p 4d 4f

3s 3p 3d

2s 2p

1s

Principal Level Distribution

Indicates the number of electrons with the sameprincipal quantum number (n). e.g.

i) Zinc: ii) Calcium +2

iii) Neon: iv) Mn:

Orbital Diagrams

Represents the electron configuration and electron spin direction using boxes to represent orbitals e.g.

i) Zinc:

ii) Calcium+2

iii) Neon:

iv) Mn:

Periodic TrendsAtomic RadiusThe distance from the center of the nucleus to the outermost electrons occupying an energy level is the atomic radius. Factors affecting atomic radii include:i) Principal Quantum Number (n)Probability of finding electrons increases farther away from the nucleus.

? as n n atomic radius n.ii) Effective Nuclear Charge (Zeff)The net force of attraction between electrons and the nucleus they surround.

? as Zeff n atomic radius p.A neutral atom will have a smaller atomic radius than the anion of the same element. e.g. Cl

A neutral atom will have a larger atomic radius than the cation of the same element. e.g. Li

Ionization Energy (kJ)

The energy required to remove an electron from an atom is called the ionization energy.

The 1st I.E. is the energy required to remove the weakest held electron from the neutral atom.

Trends in Ionization energy are linked to atomic radius and therefore are under control of:

i) Principal Quantum Number (n)? as n n I.E. p.

ii) Effective Nuclear Charge (Zeff)? as Zeff n I.E n.

Also, some variations in I.E. trends can be explained by the presence of half filled p orbitals or full s orbitals. (e.g. Be,B and N,O).

Exceptions to energy filling

¾evidence suggests that half-filled and filled subshells are more stable (lower energy) than unfilled subshells

¾i.e. Cr: [Ar] 4s2 3d4 expected but in fact it is [Ar] 4s1 3d5appears more important for d subshells and in the case of chromium an s electron is promoted to the d subshell to create two half-filled subshells

¾the justification is that the overall energy state is lower after the promotion of the electron

¾Other examples include:

Electron Affinity (kJ/mol)Electron affinity is the energy given off when an electron is added to an atom.

The 1st E.A results in the formation of an anion with a charge of 1-

Trends in electron affinity are more irregular than those for atomic radius and ionization energy.

Intermolecular Forces

Intermolecular forces act between molecules or ions. These forces must be overcome during a physical change.

i) Dipole-Dipole e.g. CHCl3

Hydrogen bonding is a special case of dipole-dipolebonding that involves a molecule containing H and a very electronegative element. e.g. H2O, NH3 etc.

ii) Ion-Dipole e.g. Na+ and H2O

iii) Ion-Induced Dipole e.g. Na+ and O2

iv) Dipole-Induced Dipole e.g. HCl and O2

v) Dispersion Forces e.g. F2 and F2

Drawing Lewis Structures for Complex Molecules

The following steps can help when determining the Lewis Structure for more complex molecules. e.g. NO3

-

1. Identify the central atom (usually lowest EN).

2. Determine the total number of Noble Gas e-

3. Determine the total number of valence e-

4. Determine the # of bonds by (taking total NG e-

subtract total valence e-) and divide by 2.

5. Determine the number of lone pair e- by subtracting the total valence e- from the number of e- involved in bonding.

6. Place lone pairs of e- around outside atoms first and then to the central atom if necessary.

Lewis Structures and Co-ordinate Covalent Bonds

A co-ordinate covalent bond is a covalent bond in which one atom donates both bonding electrons. A filled atomic orbital overlaps with an empty atomic orbital. (This is an exception to regular covalent bonding where one electron is donated by each atom.) For example;

1. Ammonium (NH4+1)

2. Hydronium (H3O+1)

3. Nitrosyl trifluoride (NF3O)

Lewis Structures and Resonance Structures

When a structure containing a double bond can be drawn with the double bond in two or more locations without changing the arrangement of atoms, a resonance structure is said to exist. These different structures for the same molecule are called “resonance hybrids”. For example:

Sulfur trioxide (SO3)

The bond lengths and strengths have been experimentally determined to fall between single and double bonds.

The electrons forming the double bond(s) are said to be “delocalized” (shared) over all the bonds.

Lewis Structures and Exceptions to the Octet Rule

1. Molecules which have more than four atoms bonded to the central atom. e.g. PCl5

2. Molecules that contain no double bonds and whose central atom has fewer than four bonding electrons. e.g. BF3

3. Molecules containing an odd number of electrons. e.g. NO

Molecules of this type are called “paramagnetic” because the are attracted by a magnetic field.

Valence Shell Electron Pair Repulsion (VSEPR) Theory

Visualize bonding and/or lone (unbonded) pairs of electrons in the outer (valence) energy level as negatively charged “clouds” (orbitals).

Each negative cloud tends to repel other negative clouds in the vicinity (around the central atom).

To achieve the most stable condition (minimum potential energy), the negative clouds must be as far apart as possible (in 3-D space).

In this position, the electrostatic repulsion is reduced to a minimum. In terms of decreasing repulsion: LP=lone pair

BP=bond pair

LP--LP > LP--BP > BP>BP

It is the position of the electron clouds around the central atom which determines the shape of the molecule. Copy figure 4.11 on page 179.

The following diagrams represent the five basic electron-group arrangements and their bond angles:

1. Linear (180O)

2. Trigonal Planar (120o)

3. Tetrahedral (109.5o or less)

4. Trigonal Bipyramidal (90o and 120o)

5. Octahedral (90o)

-tetrahedral (109.5o)

-trigonal pyramidal (107.3o)

-angular/bent/v-shape (104.5o)

-trigonal bipyramidal

-seesaw

-T-shaped

-linear

-octahedral

-square-based pyramid

-square planar

-linear

-trigonal planar

-angular/bent/v-shape